This course will cover the topics of a full year, two semester General Chemistry course. We will use a free on-line textbook, Concept Development Studies in Chemistry, available via Rice’s Connexions project. The fundamental concepts in the course will be introduced via the Concept Development Approach developed at Rice University. In this approach, we will develop the concepts you need to know from experimental observations and scientific reasoning rather than simply telling you the concepts and then asking you to simply memorize or apply them. So why use this approach? One reason is that most of us are inductive learners, meaning that we like to make specific observations and then generalize from there. Many of the most significant concepts in Chemistry are counter-intuitive. When we see where those concepts come from, we can more readily accept them, explain them, and apply them. A second reason is that scientific reasoning in general and Chemistry reasoning in particular are inductive processes. This Concept Development approach illustrates those reasoning processes. A third reason is that this is simply more interesting! The structure and reactions of matter are fascinating puzzles to be solved by observation and reasoning. It is more fun intellectually when we can solve those puzzles together, rather than simply have the answers to the riddles revealed at the outset. Recommended Background: The class can be taken by someone with no prior experience in chemistry. However, some prior familiarity with the basics of chemistry is desirable as we will cover some elements only briefly. For example, a prior high school chemistry class would be helpful. Suggested Readings: Readings will be assigned from the on-line textbook “Concept Development Studies in Chemistry”, available via Rice’s Connexions project. In addition, we will suggest readings from any of the standard textbooks in General Chemistry. A particularly good free on-line resource is Dickerson, Gray, and Haight, "Chemical Principles, 3rd Edition". Links to these two texts will be available in the Introduction module.
CDS 4 Electron Shell Model of an Atom I
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From the course by Rice University
General Chemistry: Concept Development and Application
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From the lesson
Structure of an Atom and the Electron Shell Model
Proving the existence of atoms and knowing that they combine to form molecules does not provide a means to predict how or why these atoms might combine. This requires greater detail about the structure and properties of individual atoms. In this module, we extend our understanding of atoms by making observations which reveal the internal structure of the atom including a model for the arrangement of the electrons around the atomic nucleus.
Meet the Instructors
John Steven HutchinsonDean of Undergraduates and Professor of Chemistry
In this lecture we're going to continue to understand the structure of an atom by
beginning our study of how the electrons move in the great space around the nucleus
in the atom. You'll recall that in our previous
lecture, we discussed the concept of the atomic structure, and in particular the
nuclear structure of an atom. The idea here is that there's a tiny
nucleus surrounded by vast, open space in which the electrons move.
Let's sort of try to draw a picture of what that looks like.
So we've got this very tiny, positively charged nucleus here.
It turns out that the diameter of this nucleus is o of the order of 10 to the
minus 15. To 10 to the minus 14 meters, so that's
extremely small. How is that in comparison to the atom?
Well, the radius of the atom turns out to be much larger than that.
So we have this vast empty space out here in which the electrons move, and the
radius of the atom itself turns out to be more of the order of 10 to the minus 10.
To 10 to the minus 9 meters. So you'll notice that there's a difference
in these ratios that's of the order of between 10,000 and maybe a 100,000.
Those are fairly abstract numbers. We certainly don't understand how large 10
to the minus 15, or 10 to the minus 14 meters might be.
But what we can say is some metaphor to that, or some analogy to what that might
look like. Here's a picture from the concept
development study. That might help you have an idea just how
tiny the nucleus is, relative to the size of the rest of the atom.
So here's a map of the city of Houston, and what we've done here, Rice University
is right in the middle of this map about this location here.
We've put a basketball down right in the middle of the Rice Campus here, of course
this basketball is not really drawn to scale, it should be much, much smaller,
but then we couldn't actually see it. So, imagine the size of a basketball, if
that is the size of a nucleus, then the size of the atom is about this radius that
we've drawn out here, which is about 12 kilometers large.
If you're familiar with the city of Huston, This highway here, 610, is what we
call the loop, and has a radius of around seven or eight miles, and in comparison
then, this circle is about the size of that, in comparison, the size of the
nucleus is something on the order of a basketball.
So that's actually extremely small. So one of the things we've already learned
then is that the nucleus is positively charged, but is quite tiny.
But virtually all of the mass of the atom, is inside of that tiny positively charged
mass. And the electrons are moving in this huge
open space then out here, surrounding that nucleus.
In addition, we ran an experiment in the last study where we learned something
about the atomic number, the atomic number was simply a number we assigned based upon
the mass rank of the elements. But we've now shown that the atomic number
is actually equal to the positive charge on the nucleus and since atoms are
neutral, then it's also equal to the number of electrons which are outside of
the nucleus. Now that we know that information, it
seems reasonable that we should assume that the chemistry of a particular atom
should be determined. By what these numbers of positive charges
are. That it might seem to make sense to us
since each atom has its own unique atomic number and has its own unique chemistry
that the chemistry of an atom is related to the numbers of positive charges and the
number of negative charges. Given that's the case.
Let's actually make a comparison now, of some elements that have similar atomic
numbers, and remember if they have similar atomic numbers then they have similar
numbers of positive charges in their nuclei.
And similar numbers of pos-, of, of, of negative charges, for the numbers of
electrons moving around. Here are five consecutive elements, that
we've picked out here. That have very similar atomic numbers, 8,
9, 10, 11, and 12. And therefore, similar positive charges.
What we see in this particular graph is that although they have very similar
positive charges, the elements are very different.
Oxygen and fluorine are both very reactive diatomic ga, gases.
By contrast neon, which has only one more positive charge than fluorine, is very
non-reactive, monatomic, or atomic gas. It doesn't really form any compounds at
all. And although sodium has only one more
positive charge than neon, it's a completely different material from neon.
It is a highly reactive, solid metal, and likewise magnesium is also a solid
reactive metal, although it's less reactive than sodium is.
What we can conclude from this comparison is in fact, knowing that we have similar
numbers of electrons, does not tell us that we have similar kinds of elements.
To the contrary, actually. Similar numbers of electrons might produce
very different chemical properties. But that seems contrary to what we were
thinking before that somehow or another, the number of electrons should have
something to do with the properties of the particular atom.
Let's try then to look for elements that do have similar kinds of properties.
Here's a group of them. We're going to take this particular group
here of fluorine, chlorine, bromine and iodine.
Turns out that they are quite similar. They are all diatomic elements.
Let's see. Fluorine and Chlorine are both gases at
room temperature. Bromine is a liquid at room temperature,
but it boils at a temperature not far above room temperature.
Iodine's actually a solid which sublimes into the gas phase at temperatures not far
above room temperature. So notice that these things are fairly
similar but notice also that they have very different atomic numbers.
That means they have very different positive charges and they have very
different negative charges. So we can conclude that we can get very
similar properties from very different numbers of electrons.
That tells us that the structure of the atom arrangement of electrons about the
atom must somehow or another be a bit more subtle and complicated than what we have
learned in our structure so far. So what were going to do next is to take a
look at some of the elements of the periodic table and noticed that infect
those elements can often be grouped. According to their similar properties.
That actually is what we were illustrating in the previous slide here.
This particular set of elements here turns out to be a group.
The elements fluorine, chlorine, bromine and iodine, with very similar chemical and
physical properties, they form very similar kinds of chemical compounds.
Because they are so similar, we give them a name and we call them the halogens.
There are other elements such as say sodium, potassium, rubidium, say lithium
if I list them in the order of their increasing atomic masses that are all
relatively soft highly reactive metals and we refer to these as the alkaline metals.
And of course the simplest group perhaps would be helium, neon, argon, krypton and
so forth, which we refer to as the noble gases in recognition of the fact that they
are either unreactive or extremely unlikely to form compounds.
So the first step in our process here is to recognize that elements can be grouped
according to similar properties. The next result is the very surprising
one, which is that if we look at the elements in the groups, for example this
fluorine, chlorine, bromine, or iodine. They appear in the Periodic Table
periodically. So let's actually pull up a table here
that illustrates that that's the case. This is actually from the concept
development study on the shell model of an atom.
So let's go look for elements that have similar kinds of properties here.
Here, for example, is fluorine. And here, for example, is chlorine.
And, let's see, slide this up a little up further so that we can see it, and down
below, we'll actually find bromine. And if we notice where these elements are,
in each case they, the elements in the group which are called halogens
immediately precede the elements which we call the noble gases.
Notice fluorine precedes neon, chlorine precedes argon and bromine precedes
krypton down below. And after each noble gas, including if we
look up, back above to helium. Immediately after helium is a member of
the alkaline metals, and immediately after neon is a member of the alkaline metals.
And immediately after argon is a member of the alkaline metals, and I'll look down at
the bottom, we have not shown this, the next element down below is another
alkaline metal, called rubidium. The point here is that the different
groups of elements such as for example lithium, and sodium, and potassium appear
recurrently, they appear periodically in the periodic table.
That actually is an example of something we call the periodic law.
The periodic law says. That if we rank the elements by their
atomic number the physical and chemical properties of the elements are all
periodic functions of that atomic number. That is, the properties recur over and
over again as we go through the list order of the elements.
That causes to give rise to something called the periodic table of the elements.
You can get access to a periodic table in lots of different locations, but it's a
beautiful construct that allows us to look at, the relationship of the different
elements to each other. This particular one is taken from the
Royal Society of Chemistry's website. They call it their visual elements
periodic table. Notice in the periodic table here, what
we've done is to list the periodic elements moving across the periodic table
here as we increase the atomic number. But we've also stacked up the elements
which have similar chemical and physical properties.
So fluorine, chlorine, bromine, and iodine are the four elements we listed as
halogens. Helium, neon, argon, krypton, xenon, are
the elements we listed as noble gases. Lithium, sodium, potassium, rubidium are
the elements that we listed as the alkaline metals.
In each case, each column of the periodic table corresponds to a set of elements
which have somewhat similar chemical and physical properties.
Whereas the rows are simply increasing numbers of, of atomic numbers, the
increasing number in the mass order. Let's go back to our slides here then.
And say that we now discovered that although the, it's not a simple
relationship between the number of electrons and the number of protons in an
atom, it appears to be a periodic function of the number of electrons and protons in
an atom which determines its chemical and physical properties.
We're going to consider one very important example of the elements and their
periodicity which is something called ionization energy.
The ionization energy as expressed here it is the amount of energy required to take
an electron away from an atom. Remember that our atom looks something
like this. Positively charged nucleus surrounded in
some large space here in some way or another with electrons flying around.
And the idea here is what if I were to try to remove that electron.
If I took the electron away what would remain behind would be a positive ion.
But it would cost us energy to be able to pull that electron away from the atom and
that's a consequence of the fact that there's a force of attraction between the
electron and the positively charged nucleus, very much like these two magnets
that I've got here. These are magnets you get in the toy
store. In this case, if I want to pull these two
magnets apart as they are attracted to each other, I have to add energy to them
to pull them apart. I have to do some work.
Similarly if I let the atoms or the magnets come back together, actually
energy is released as they attract each other.
So, energy is absorbed when I pull them apart and is released when I bring them
back together. Therefore, if I want to remove an electron
from an atom I am going to have to apply an ionization energy.
That is the minimum amount of energy which is required to remove an electron.
Without going into details about how we are going to measure the ionization energy
let's actually just look at what the data appear to be as we look at the ionization
energies. Here's a graph of the minimum energy
required to remove an electron from each of the elements of the periodic table and
we've for a set of them and we've simply graphed them as a function now of the
atomic number. And notice that you actually see the
groups in the periodic table in kind of conspicuous and similar sorts of
locations. So, for example, right at the top of each
of these, we see the noble gases, and right at the bottom of the chart in each
case, we see the alkali metals. So there appears to be groups that have
similar kinds of properties so the periodic law is really quite conspicuous
here. Well what's going on in this graph in
addition? Let's remove these markings here to make
them a little easier to see. Let's go back and look at our list of
elements. Or actually wanted to look at the periodic
table here. And bring it across to where we can
compare the periodic table to the ionization energies here.
Notice in each case that as we move across the periodic table from, say.
Lithium, beryllium, boron, carbon, all the way out to neon.
That, that corresponds to this particular stretch moving from lithium to neon.
And then if we go to the next row of the periodic table, sodium, magnesium, going
all the way across to argon, that corresponds to this rise.
And notice we keep getting these periodic rises as we move from the alkali metals
all the way across to the noble gases. Within an addition we get this rather
precipitous drop each time we move from a noble gas down to the subsequent alkali
metal even back over here as well. So there are some clear repeating patterns
in this data that have something to do with the arrangement of the electrons in
the table. For now we're simply going to draw two
conclusions. The first is that when you measure the
ionization energies of the atoms, the periodic law is really quite clear in this
state. This actually looks like a periodic
function. And further more we can easily see the
appearance of the groups of elements we described earlier in this lecture
appearing in this data. Why the ionization energy looks this way
what we're going to take up in the next lecture when we start to understand
exactly how this data tells us the arrangement of the electrons about the
nucleus.
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